Notes on Atoms, Molecules, and Ions (Chapter 2)
Atomic Theory and Early Discoveries
Atomic Theory of Matter reemerged in the early nineteenth century, championed by John Dalton.
Dalton’s Postulates:
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical to one another in mass and other properties, but atoms of one element differ from atoms of all other elements.
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
Atoms of more than one element combine to form compounds; a given compound always has the same relative number and kind of atoms.
Law of Conservation of Mass:
The total mass before a chemical process equals the total mass after the process. m{ ext{before}} = m{ ext{after}}
This law underpinned Dalton’s atomic theory.
Law of Multiple Proportions:
If two elements, A and B, form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers. Dalton predicted this law and observed it while developing his atomic theory.
When two or more compounds exist from the same elements, they cannot have the same relative number of atoms.
Subatomic Particles and the Electron
Discovery of Subatomic Particles:
Dalton’s atom viewed as the smallest particle, but experiments showed atoms contain smaller particles: electrons, cathode rays, radioactivity, nucleus, protons, and neutrons.
The Electron (Cathode Rays):
Streams of negatively charged particles emanate from cathode tubes, causing fluorescence.
J. J. Thomson credited with discovery (1897).
Electron Charge-to-Mass Ratio:
Thomson measured the charge/mass ratio of the electron to be rac{e}{m_e} = 1.76 \times 10^{8} \; ext{C g}^{-1}
Millikan Oil-Drop Experiment (Electrons):
Once the charge/mass ratio of the electron was known, determining either the charge or the mass yielded the other.
Robert Millikan determined the charge on the electron in 1909.
Radioactivity and Nuclear Model
Radioactivity:
The spontaneous emission of high-energy radiation by an atom.
First observed by Henri Becquerel; studied by Marie and Pierre Curie.
Its discovery showed that the atom has more subatomic particles and energy associated with it.
Types of Radiation (Ernest Rutherford):
alpha (a) particles – positively charged
beta (b) particles – negatively charged, like electrons
gamma (g) rays – uncharged
The Atom at the Turn of the 20th Century
The Plum Pudding Model (Thomson):
A positive sphere of matter with embedded negative electrons.
Discovery of the Nucleus (Rutherford):
Shot a-particles at a thin sheet of gold foil and observed the scattering pattern.
Some particles were deflected at large angles, inconsistent with Thomson’s model.
The Nuclear Atom (Rutherford):
A very small, dense nucleus containing protons and neutrons, with electrons around the outside.
Most of the atom’s volume is empty space.
Atoms are very small; typical sizes are about 1$-$5 \text{ \AA} or 100$-$500 \text{ pm}.
Protons and neutrons were discovered as subatomic particles within the nucleus; electrons travel around the nucleus.
Subatomic Particles and Atomic Structure
Subatomic Particles:
Protons (+1) and electrons (−1) have a charge; neutrons are neutral.
Protons and neutrons have essentially the same mass (relative mass 1). The mass of an electron is so small we ignore it (relative mass 0).
Protons and neutrons are in the nucleus; electrons travel around the nucleus.
Atomic Mass and Mass Units:
Atoms have extremely small masses; the heaviest known atoms have a mass of approximately 4 \times 10^{-22} \text{ g}.
An atomic mass unit (amu) is the base unit for atomic-scale masses: 1 \text{ amu} = 1.66054 \times 10^{-24} \text{ g}.
Atomic Weight and Measurement:
Atomic and molecular weights can be measured with great accuracy using a mass spectrometer.
Masses of atoms are compared to the carbon atom with 6 protons and 6 neutrons (C-12).
Isotopes and Atomic Weight
Symbols of Elements:
Elements are represented by a one- or two-letter symbol.
Atomic number Z is the number of protons; it is written as a subscript before the symbol.
Mass number A is the total number of protons and neutrons in the nucleus; it is written as a superscript before the symbol.
Isotopes:
Isotopes are atoms of the same element with different masses.
Different numbers of neutrons, but the same number of protons.
Atomic Weight (Average Mass):
Because large amounts of atoms are used, we use average masses weighted by isotope abundances.
Atomic Weight = \sum (\text{isotope mass}) \times (\text{fractional natural abundance})
The sum runs over ALL isotopes of the element.
Periodic Table Basics
Periodic Table Organization:
Elements arranged in order of atomic number.
The atomic number (Z) is at the top of each box; the atomic weight appears at the bottom (in this version not always shown).
Structure:
Rows = periods; Columns = groups.
Elements in the same group have similar chemical properties.
Periodicity:
Observed repeating patterns in chemical properties as you move across periods and groups.
Classification:
Metals are on the left side of the table; some properties include shine, conductivity of heat and electricity, and solidity (except mercury).
Nonmetals are on the right side (except H); states at room temperature vary: solid (C), liquid (Br), gas (Ne).
Metalloids lie on a steplike line (except Al, Po, At), with properties between metals and nonmetals.
Groups and Names in the Periodic Table
Group names (examples):
1A: Alkali metals — Li, Na, K, Rb, Cs, Fr
2A: Alkaline earth metals — Be, Mg, Ca, Sr, Ba, Ra
6A: Chalcogens — O, S, Se, Te, Po
7A: Halogens — F, Cl, Br, I, At
8A: Noble gases — He, Ne, Ar, Kr, Xe, Rn
The five groups are commonly named as listed above.
Formulas, Ions, and Ionic Compounds
Formulas and Molecules:
The subscript to the right of an element’s symbol indicates the number of atoms of that element in one molecule of a compound.
Molecular compounds are composed of molecules and almost always contain only nonmetals.
Diatomic Molecules:
Seven elements naturally occur as diatomic molecules: \mathrm{H2, N2, O2, F2, Cl2, Br2, I_2}
Types of Formulas:
Empirical formulas give the lowest whole-number ratio of atoms in a compound.
Molecular formulas give the actual number of atoms of each element in a compound.
If the molecular formula is known, the empirical formula can be determined; the converse is not always true.
Structural and Perspective Formulas:
Structural formulas show the order in which atoms are attached.
Perspective drawings show the three-dimensional order of atoms; models also illustrate these arrangements.
Ions:
Atoms (or groups) that lose or gain electrons become ions.
Cations form when at least one electron is lost; monatomic cations come from metals.
Anions form when at least one electron is gained; monatomic anions come from nonmetals.
Common Ions and Nomenclature
Common Cations (selected):
1+: H+, Li+, Na+, K+, Cs+; NH4+ (ammonium); Cu+ (copper(I))
2+: Mg2+, Ca2+, Sr2+, Ba2+, Zn2+, Ni2+, Fe2+ (ferrous), Co2+ (cobaltous), Cu2+ (cupric), Hg2^2+ (mercurous/mercuric as applicable), Pb2+ (lead(II)), Sn2+ (tin(II))
3+: Al3+, Cr3+, Fe3+ (ferric)
Common Anions (selected):
1−: F−, Cl−, Br−, I−, OH−, CN−, CH3COO− (acetate, or C2H3O2−), NO3− (nitrate), NO2− (nitrite), ClO− (hypochlorite), ClO3− (chlorate), ClO4− (perchlorate), MnO4− (permanganate)
2−: O2− (oxide), CO3^2− (carbonate), SO3^2− (sulfite), SO4^2− (sulfate), CrO4^2− (chromate), Cr2O7^2− (dichromate), O2^2− (peroxide)
3−: PO4^3− (phosphate)
Ionic Compounds:
Ionic compounds generally form between metals and nonmetals.
Electrons are transferred from the metal (cation) to the nonmetal (anion); the oppositely charged ions attract.
Only empirical formulas are written for ionic compounds.
Writing Formulas for Ionic Compounds:
The charge on the cation becomes the subscript on the anion, and the charge on the anion becomes the subscript on the cation.
If the subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.
Example guidance: NaCl (sodium chloride) from Na+ and Cl−; CaCl2 from Ca2+ and Cl−; etc.
Inorganic Nomenclature and Oxyanions
Naming Cations and Anions:
Name the cation first; if the cation can have more than one charge, indicate the charge with a Roman numeral in parentheses.
If the anion is a simple element, change its ending to -ide; if it is a polyatomic ion, use its full name.
Oxyanion Nomenclature Patterns:
Two oxyanions of the same element: the one with fewer oxygens ends in -ite; the one with more oxygens ends in -ate.
Examples: NO2− (nitrite) vs NO3− (nitrate); SO3^2− (sulfite) vs SO4^2− (sulfate).
Central atoms on the second row tend to bond to at most three oxygens; third-row central atoms can bond to four.
When there are two oxyanions with different oxygen counts: the one with the second fewest oxygens ends in -ite (e.g., ClO2− is chlorite); the one with the second most ends in -ate (e.g., ClO3− is chlorate).
The fewest-oxygen ion has the prefix hypo- and ends in -ite (e.g., ClO− is hypochlorite); the most-oxygen ion has the prefix per- and ends in -ate (e.g., ClO4− is perchlorate).
Acid Nomenclature:
If the anion ends in -ide, the acid name is hydro- + -ic acid (e.g., HCl: hydrochloric acid; HBr: hydrobromic acid; HI: hydroiodic acid).
If the anion ends in -ite, the acid ends in -ous acid (e.g., HClO: hypochlorous acid; HClO2: chlorous acid).
If the anion ends in -ate, the acid ends in -ic acid (e.g., HClO3: chloric acid; HClO4: perchloric acid).
Binary Molecular Compounds (nonmetals):
The element farther to the left (or lower in the same group) is named first.
A prefix indicates the number of atoms of each element (mono- is not used for the first element).
The ending of the second element is -ide (e.g., CO2: carbon dioxide; CCl4: carbon tetrachloride).
If prefixes create vowel adjacency, vowels may be elided (e.g., N2O5: dinitrogen pentoxide).
Organic Nomenclature
Organic chemistry studies carbon; it has its own nomenclature system.
The simplest hydrocarbons are alkanes (contain only C and H).
The number of carbons is indicated by the root: meth- (1), eth- (2), prop- (3), etc.
When a hydrogen in an alkane is replaced with a functional group (e.g., -OH in alcohols), the name is derived from the alkane name.
The ending denotes the type of compound; for alcohols, the suffix is -ol.
Summary of Key Formulas and Conventions
Atomic mass unit: 1 \text{ amu} = 1.66054 \times 10^{-24} \text{ g}
Electron charge-to-mass ratio: \frac{e}{m_e} = 1.76 \times 10^{8} \ \text{C g}^{-1}
Isotope notation: ^{A}_{Z}\text{X} where A = mass number, Z = atomic number, X = element symbol, and A = Z + N with N neutrons.
Atomic Weight: =\Sigma\left(mass\cdot abundance\right) of each isotope
Diatomic elements: Iodine, Hydrogen, Nitrogen, Bromine, Oxygen, Chlorine, Fluorine
Ion charges and formula balance concepts are described in the balance rules above.
Structural vs perspective naming and organic nomenclature basics are described in the sections above.
Connections to Foundational Principles and Real-World Relevance
Dalton’s postulates laid the groundwork for chemical reactions, stoichiometry, and conservation laws used in laboratory experiments and industrial processes.
The conservation of mass underpins all chemical equations and quantitative chemistry.
The discovery of subatomic particles and the nuclear model revolutionized our understanding of matter, leading to modern physics, nuclear chemistry, and applications such as radiometric dating, medical imaging, and energy generation.
Periodic table organization reflects recurring chemical properties, enabling prediction of element behavior and compound formation.
Isotopes and atomic weight explain natural variations in element masses and are crucial for accurate chemical calculations and applications in spectroscopy, pharmacology, and materials science.
Nomenclature rules (ionic, oxyanion, acid, binary molecular, and organic) provide a universal language for communicating chemical formulas and structures, essential for science, medicine, and industry.